Atomic Structure and Bonding

🟢 Lite — Quick Review (1h–1d)

Rapid summary of atomic structure and bonding for NABTEB chemistry.

The Atom is the basic unit of matter. Key subatomic particles:

Particle	Symbol	Charge	Mass (approx)	Location
Proton	$p^+$	+1	1 u	Nucleus
Neutron	$n^0$	0	1 u	Nucleus
Electron	$e^-$	-1	1/1836 u	Electron shells

Atomic Number ($Z$): Number of protons in the nucleus. Unique to each element. Mass Number ($A$): Total number of protons + neutrons in the nucleus. Isotopes: Atoms of the same element with different mass numbers (different number of neutrons).

Electron Configuration: Electrons occupy shells (energy levels) around the nucleus. Each shell holds a maximum number of electrons:

• Shell 1 (K): maximum 2 electrons
• Shell 2 (L): maximum 8 electrons
• Shell 3 (M): maximum 18 electrons (but 8 for first 20 elements)
• Shell 4 (N): maximum 32 electrons

Filling Order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p…

Lewis Structure:

• Dots represent valence electrons (electrons in outer shell)
• Atoms bond to achieve 8 electrons in outer shell (octet rule), except H (2 electrons).

⚡ NABTEB Exam Tip: The number of valence electrons = Group number (for Groups 1-18). Group 1 elements have 1 valence electron, Group 2 have 2, Group 13 have 3, and so on.

---

🟡 Standard — Regular Study (2d–2mo)

For NABTEB students who want solid understanding of atomic structure and bonding.

Types of Chemical Bonds:

1. Ionic Bonding:

• Transfer of electrons from metal to non-metal
• Creates oppositely charged ions held by electrostatic attraction
• Example: $NaCl$ — Na loses 1 electron → $Na^+$; Cl gains 1 electron → $Cl^-$
• Properties: High melting/boiling points, conduct electricity when molten or dissolved, often soluble in water

2. Covalent Bonding:

• Sharing of electron pairs between non-metal atoms
• Can be single ($H_2$), double ($O_2$), or triple ($N_2$) covalent bonds
• Can be polar (unequal sharing, e.g., $HCl$) or non-polar (equal sharing, e.g., $CH_4$)
• Properties: Low melting/boiling points, do not conduct electricity (except graphite)

3. Metallic Bonding:

• Metal ions in a sea of delocalised electrons
• Explains malleability, conductivity, and lustre of metals
• Properties: High melting points, conduct electricity in solid and liquid states

Electronegativity:

A measure of the tendency of an atom to attract a bonding pair of electrons.

• Values range from 0.7 (Cs) to 4.0 (F)
• Difference in electronegativity predicts bond type:

  • 1.7: Ionic bond

  • 0.4–1.7: Polar covalent
  • < 0.4: Non-polar covalent

Common Electronegativity Values:

Element	Electronegativity
F	4.0
O	3.5
N	3.0
Cl	3.2
C	2.5
H	2.1
Na	0.9
K	0.8

Bonding and Structure:

Substance	Bonding Type	Structure	Properties
NaCl	Ionic	Giant lattice	Brittle, high melting, conducts when molten
Diamond	Covalent (network)	Giant covalent	Very hard, high melting, non-conductor
Graphite	Covalent	Layered	Soft, slippery, conducts electricity
$CO_2$	Covalent (simple)	Simple molecular	Low melting, non-conductor
Iron	Metallic	Metallic lattice	Malleable, conducts electricity

Octet Rule Exceptions:

• Hydrogen: achieves 2 electrons (duet)
• Boron: stable with 6 electrons ($BF_3$ is electron-deficient)
• Phosphorus: can expand octet ($PCl_5$, 10 electrons)
• Sulphur: can expand octet ($SF_6$, 12 electrons)
• Noble gases: complete octet already (Group 18)

⚡ NABTEB Exam Tip: Ionic compounds conduct electricity when molten or in solution because ions become mobile. Covalent compounds never conduct electricity through solid or liquid states (unless they ionise in water, like HCl).

---

🔴 Extended — Deep Study (3mo+)

Comprehensive coverage of atomic structure and bonding for thorough NABTEB preparation.

Quantum Numbers:

Each electron in an atom is described by four quantum numbers:

1. Principal quantum number ($n$): Shell/energy level (1, 2, 3, 4…)
2. Orbital angular momentum ($l$): Sub-shell type — $s$ (0), $p$ (1), $d$ (2), $f$ (3)
3. Magnetic quantum number ($m_l$): Orientation of orbital (−l to +l)
4. Spin quantum number ($m_s$): Direction of spin — $+\frac{1}{2}$ or $-\frac{1}{2}$

Electron Configurations:

Using $nl^x$ notation:

• $1s^2$: 2 electrons in 1s orbital
• $2s^2 2p^6$: Full second shell (8 electrons)
• $3s^2 3p^3$: Phosphorus — has 5 valence electrons

Orbital Filling Rules:

1. Aufbau principle: Fill lowest energy orbitals first
2. Pauli exclusion principle: Max 2 electrons per orbital (opposite spins)
3. Hund’s rule: Fill orbitals singly before pairing (maximises unpaired electrons)

Orbital Shapes:

• s orbital: Spherical (1 per shell)
• p orbital: Dumbbell-shaped (3 per sub-shell: $p_x, p_y, p_z$)
• d orbital: Double-dumbbell/cloverleaf (5 per sub-shell)
• f orbital: Complex shapes (7 per sub-shell)

Bonding Molecular Orbitals:

In covalent bonding:

• Bonding orbitals: Lower energy, electrons prefer to occupy
• Antibonding orbitals: Higher energy, electrons fill after bonding orbitals are full
• Bond order = (Number of bonding electrons − Number of antibonding electrons) ÷ 2
• Bond order > 0 means stable molecule

Sigma ($\sigma$) and Pi ($\pi$) Bonds:

• $\sigma$ bond: Head-on overlap of orbitals; strongest; present in all bonds
• $\pi$ bond: Sideways overlap of p orbitals; weaker; found in double and triple bonds

Bond Type	Number of $\sigma$ bonds	Number of $\pi$ bonds
Single (C–C)	1	0
Double (C=C)	1	1
Triple (C≡C)	1	2

VSEPR Theory (Valence Shell Electron Pair Repulsion):

Electron pairs around a central atom repel each other and arrange to minimise repulsion:

Steric Number	Shape	Bond Angles
2	Linear	180°
3	Trigonal planar	120°
4	Tetrahedral	109.5°
5	Trigonal bipyramidal	90°, 120°
6	Octahedral	90°

Lone pairs repel more strongly than bonding pairs, so bond angles are reduced.

Hybridisation:

Hybridisation	Orbitals Mixed	Geometry	Examples
$sp$	one s + one p	Linear	$BeCl_2$, $C_2H_2$
$sp^2$	one s + two p	Trigonal planar	$BF_3$, $C_2H_4$
$sp^3$	one s + three p	Tetrahedral	$CH_4$, $C_2H_6$
$sp^3d$	one s + three p + one d	Trigonal bipyramidal	$PCl_5$
$sp^3d^2$	one s + three p + two d	Octahedral	$SF_6$

Intermolecular Forces:

Forces between molecules (weaker than covalent bonds):

1. Van der Waals (London dispersion): Temporary dipoles in all molecules; stronger with more electrons
2. Dipole-dipole: Permanent dipoles in polar molecules; e.g., HCl
3. Hydrogen bonding: Strong dipole-dipole when H is bonded to F, O, or N (N, O, F have high electronegativity)

Substance	Intermolecular Force	Boiling Point
$He$	London	-269°C
$H_2$	London	-253°C
$N_2$	London	-196°C
$O_2$	London	-183°C
$HCl$	Dipole-dipole	-85°C
$NH_3$	Hydrogen bonding	-33°C
$H_2O$	Hydrogen bonding	100°C

Properties of Water — Why it boils at 100°C:

Water has strong hydrogen bonding between molecules, requiring significant energy to overcome. This explains its relatively high boiling point for its molecular mass.

Ionic vs Covalent Character:

Electronegativity difference predicts bond character:

• Difference > 1.7: Predominantly ionic
• Difference 0.4–1.7: Polar covalent
• Difference < 0.4: Predominantly non-polar covalent

⚡ NABTEB Quick Reference:

• $A = Z + N$ (mass number = protons + neutrons)
• Aufbau: $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d$
• Ionic: metal + non-metal → electron transfer
• Covalent: non-metal + non-metal → electron sharing
• Metallic: metal + metal → electron sea
• Electronegativity difference > 1.7: ionic
• VSEPR: electron pairs repel, minimise angles
• $sp^3$: 4 bonds = tetrahedral; $sp^2$: 3 bonds = trigonal planar; $sp$: 2 bonds = linear
• Hydrogen bonding: H bonded to F, O, or N